Flashcards · 9620 · AS Topic 3

Bonding (AS Level) flashcards

32 cards on AS Topic 3, subtopics 3.1.3.1 to 3.1.3.7. Pick a subtopic or study them all.

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Ionic bonding: The electrostatic attraction between oppositely charged ions in a lattice. It acts in all directions.
Formulae and charges of the five compound ions to learn: Sulfate SO₄²⁻, carbonate CO₃²⁻, nitrate NO₃⁻, hydroxide OH⁻, ammonium NH₄⁺.
Charges on simple ions of Groups 1, 2, 3, 5, 6 and 7: Group 1: 1+. Group 2: 2+. Group 3: 3+. Group 5: 3−. Group 6: 2−. Group 7: 1−.
Formulae of aluminium sulfate and magnesium nitride: Al₂(SO₄)₃ and Mg₃N₂.

Single covalent bond: One shared pair of electrons. A double bond has two shared pairs; a triple bond has three.
Co-ordinate (dative covalent) bond: A shared pair of electrons where both electrons come from one atom. Drawn as an arrow from the donor atom to the acceptor.
What does an atom need to donate, and to accept, a co-ordinate bond?: Donor: a lone pair of electrons. Acceptor: an empty orbital (a gap in its outer shell).
Name the donor and acceptor in NH₄⁺ and in H₃N→BF₃: NH₄⁺: N of NH₃ donates to H⁺. H₃N→BF₃: N of NH₃ donates to B of BF₃.

Metallic bonding: The attraction between a lattice of positive ions and delocalised electrons.
What makes metallic bonding stronger?: A larger charge on the ions, more delocalised electrons per atom and smaller ions.
Why does magnesium have a higher melting point than sodium?: Mg forms 2+ ions and releases two delocalised electrons per atom (Na: 1+, one electron), so the attraction between ions and delocalised electrons is stronger.

The four types of crystal structure, with the six named examples: Ionic: sodium chloride. Metallic: magnesium. Macromolecular (giant covalent): diamond, graphite. Molecular: iodine, ice.
Electrical conductivity of sodium chloride: Does not conduct as a solid (ions fixed in the lattice). Conducts when molten or dissolved, because the ions are free to move.
Diamond vs graphite: bonding and conductivity: Diamond: each C bonded to 4 others in a 3-D network; no free charge carriers, so does not conduct. Graphite: each C bonded to 3 others in layers; one delocalised electron per atom, so conducts along the layers.
What is overcome when a molecular solid such as iodine melts or sublimes?: Only the weak intermolecular forces between molecules. The covalent bonds inside each molecule stay intact.
Why do melting and boiling need energy?: They are endothermic: particles are pulled apart against the forces of attraction between them (ionic bonds, metallic bonding, covalent bonds or intermolecular forces, depending on the structure).

Order of strength of repulsion between electron pairs: Lone pair–lone pair > lone pair–bond pair > bond pair–bond pair. So each lone pair reduces the bond angles.
Method to find the number of electron pairs round a central atom: Central atom’s outer electrons + one per bonded monovalent atom − the charge (add for a negative ion). Divide by two for total pairs; subtract bonding pairs to get lone pairs.
Shapes and bond angles of NH₃ and H₂O: NH₃: pyramidal, 3 bonding + 1 lone pair, about 107°. H₂O: bent, 2 bonding + 2 lone pairs, about 104.5°.
Shapes and angles: BF₃, CH₄ (or NH₄⁺), PCl₅, SF₆: BF₃: trigonal planar, 120°. NH₄⁺: tetrahedral, 109.5°. PCl₅: trigonal bipyramidal, 90° and 120°. SF₆: octahedral, 90°.
Shape of ICl₄⁻: 7 + 4 + 1 = 12 electrons, so 6 pairs: 4 bonding and 2 lone. Square planar, 90°.
Shapes of SF₄ and ClF₃: SF₄: see-saw (4 bonding, 1 lone pair). ClF₃: T-shaped (3 bonding, 2 lone pairs). In a trigonal bipyramid the lone pairs take equatorial positions.

Electronegativity: The power of an atom to attract the pair of electrons in a covalent bond towards itself.
When is a covalent bond polar?: When the two atoms have different electronegativities, so the electron density is unsymmetrical. The more electronegative atom is δ−.
Why does CCl₄ have no permanent dipole although its bonds are polar?: Four identical polar bonds are arranged symmetrically (tetrahedral), so the bond dipoles cancel.
Which of SO₃, SF₆, PCl₃ and CH₂Cl₂ have a permanent dipole?: PCl₃ (pyramidal) and CH₂Cl₂ (unlike atoms round C). In SO₃ and SF₆ the dipoles cancel.

Induced dipole–dipole forces (van der Waals, London, dispersion): Act between all molecules. Electron movement creates a temporary dipole that induces a dipole in a neighbouring molecule. Stronger with more electrons and a larger surface area of contact.
Permanent dipole–dipole forces: Attraction between the δ+ end of one polar molecule and the δ− end of another.
Conditions for hydrogen bonding: A hydrogen atom bonded to N, O or F, attracted to a lone pair on an N, O or F atom of another molecule.
Why does pentane boil higher than 2,2-dimethylpropane?: The straight chain has a larger surface area of contact with neighbouring molecules, so the induced dipole–dipole forces are stronger.
Why is ice less dense than liquid water?: In ice each molecule is held by four hydrogen bonds in an open lattice, so the molecules are further apart than in the liquid.
Why do H₂O, NH₃ and HF have anomalously high boiling points?: Hydrogen bonding between their molecules is stronger than the intermolecular forces in the hydrides of their group neighbours, so more energy is needed to separate them.